Chlorine dioxide enjoys considerable commercial and industrial importance in a wide variety of applications. It is currently used in large quantities as a bleaching agent for wood pulp, paper, fats, oils, tallow, and flour.
A recent series of regulatory approvals have increased the use of chlorine dioxide as a disinfectant and sanitizer in the food processing industry. New federal guidelines have permitted its use in meat, dairy, poultry, fruit and vegetable post-harvest produce, and prepared foodstuffs. In most circumstances, chlorine dioxide does not cause organoleptic impairments of food products.
Chlorine dioxide is also widely used in the wastewater industry both as a pollution control agent and a potable water treatment. Chlorine dioxide is an excellent sulfide scavenging agent and is employed in scrubbing towers in wastewater, rendering, and the oil and gas industry.
The unique properties of chlorine dioxide provide a growing receptivity in its use as an environmental and microbial control agent. Chlorine dioxide reacts with a high degree of specificity towards certain industrial pollutants, such as sulfides, amines, mercaptans, and cyanide while not reacting with ammonia or most organic compounds.
The highly selective nature of chlorine dioxide is important in disinfection. Unwanted disinfection by-products such as trihalomethanes (THMs) and polychlorobiphenyls (PCBs) are not formed as they are with chlorine or hypochlorites. Chlorine dioxide also is effective over a wide pH range, does not disassociate in solutions, has rapid disinfection kinetics, and does not accumulate in treated solutions.
Chlorine dioxide is most often generated on-site due to prohibition and hazards of its transport and storage. It has been produced conventionally by several chemical and electrochemical processes. The most common means of production is the acidification of aqueous sodium chlorite. Strong acids such as sulfuric or hydrochloric give high yields of chlorine dioxide while weaker (and safer) acids such as citric and lactic give much lower yields.
The acidic conversion of chlorite to chlorine dioxide is greatly enhanced in yield by adding a chlorine donor such as hypochlorite or chlorine gas. While advantageous in recoverable yield, a three-precursor system introduces a greater level of complexity in apparatus and reactor design. Current methods for producing chlorine dioxide gas employ highly toxic chlorine gas. On-site storage of hypochlorites, strong acids, and chlorine gas also poses additional hazards and regulatory scrutiny to the end user.
Electrochemical generators have attempted to partially address this issue. Such generators utilize a singe precursor, normally sodium chlorite or sodium chlorate. The produced chlorine dioxide product is separated from the electrolyte solution using a gas permeable structure. However, they are not widely used due to other disadvantages such as high cost, generation of explosive hydrogen gas, and reliability. New designs have begun to address these concerns
A potentially superior and alternative method for producing chlorine dioxide is by photochemical oxidation. Photochemical reactions of chlorine dioxide and oxyanions of chlorine have been reported by E. J. Brown and M. Cheung (1932) and disclosed in U.S. Pat. Nos. 2,043,284, 2,457,285, and 2,683,651.
More recent work of photochemical methods is disclosed in U.S. Pat. Nos. 4,414,180 and 4,456,511 to Fisher. This work describes a generator containing aqueous sodium chlorite illuminated by externally placed incandescent fluorescent bulbs. The sodium chlorite is photochemically oxidized to chlorine dioxide and removed from the aqueous solution with a gaseous nitrogen or air stream.
Further, more detailed work is disclosed in U.S. Pat. No. 4,874,489 to J. Callerame, which describes a tubular chamber containing sodium chlorite. An ultraviolet source, namely a low pressure mercury vapor bulb, was housed inside the vessel and the inside wall was made from a UV reflector such as polished aluminum. The reaction was discontinued when the chlorine dioxide concentration reached ten percent weight and the entire reaction product was removed from the reaction space. Ten percent was the upper limit chosen due to explosive properties of chlorine dioxide.
Similarly in U.S. Pat. No. 4,877,500 to Callerame, mixtures of chlorine and oxygen gas and aqueous solutions of sodium hypochlorite were photochemically converted to chlorine dioxide. As in the earlier Callerame patent, the solution was held within the tubular vessel until the maximum concentration of chlorine dioxide reached ten percent. Afterwards, the entire reaction contents containing chlorine dioxide were removed and conveyed to their place of use. In the photochemical reactions of chlorine and oxygen, explosions were reported to have occurred in two instances.
Improvements in safety of photochemical methods were reported by Simpson in U.S. Pat. No. 6,171,558. A means is described for positioning a UV bulb in a container of aqueous chlorite. The aqueous chlorite is circulated through a circulation tube by an air or gas sparge. This also effectively removes chlorine dioxide and thus reduces the safety hazardous associated with its accumulation in solution.
While considerably improving upon the safety of the prior photochemical methods the following limitations are still imposed:                (1) The prior art requires additional air moving system to generate a gas sparge.        (2) The prior art was conducted in fragile quartz tubes and aspirators.        (3) The prior art requires a circulation tube in proximity to the ultraviolet bulb to conduct the chlorite precursor across the field of ultraviolet radiation.        (4) The prior art production of chlorine dioxide cannot be directly regulated by controlled addition of chlorite precursor. The prior art production of chlorine dioxide must be controlled by bulb intensity and gas sparge rate.        (5) The scale-up of the prior art device is difficult. The chlorine dioxide production rate also decreases as the chlorite is exhausted        
Chlorine dioxide has the potential for increased use in a variety of commercial and industrial applications. The apparatus for its generation would ideally be safe, economical, easy-to-use, and not require storage of hazardous ingredients. A clear and compelling need exists for a device that produces chlorine dioxide while fulfilling the above criteria.